Calcium Sulfate

Author: Hans Lohninger

Calcium sulfate, CaSO₄, is a common laboratory and industrial chemical and an often used material in the building trade. It occurs naturally in various forms, which differ in their crystal water content: gypsum (CaSO₄·2H₂O), the hemihydrate (CaSO₄·0.5H₂O) also known as plaster of Paris, and the anhydrite. Despite its name the anhydrite is not always entirely devoid of water, the water content ranges from 0.0 to 0.05 mol-percent.

The main sources of calcium sulfate are naturally occurring minerals (gypsum and anhydrite). World production of natural gypsum is about 100 million tonnes per annum. Besides the natural sources, calcium sulfate is also produced as a by-product, mainly from the desulfurization of exhaust gases of fossil-fuel power stations.

Properties

Gypsum, CaSO₄·2H₂O, is thermodynamically stable in aqueous solutions below 42°C, while anhydrite, CaSO₄, is stable above 42°. Thus gypsum crystallizes in aqueous solutions at low temperatures, while anhydrite crystallizes at higher temperatures. The hemihydrate is meta-stable.

Alkali and alkaline earth elements form double and triple salts with calcium sulfate, e.g. CaSO₄K₂SO₄H₂O, or 2CaSO₄MgSO₄K₂SO₄2H₂O

Dehydration of Gypsum

When heating gypsum to between 80 and 180°C it loses increasing amounts of water, forming first the hemihydrate and later the anhydrite:

CaSO₄·2H₂O + heat CaSO₄·½H₂O + 1½H₂O (steam)
CaSO₄·½H₂O + heat CaSO₄ + ½H₂O (steam)

When mixed with water at ambient temperatures, calcined gypsum quickly reverts to the preferred dihydrate form, developing an interwoven network of gypsum crystals. This re-hydration reaction is exothermic and is responsible for the ease with which gypsum can be cast into various shapes.